Molarity Calculator
Calculate solution molarity, required mass, moles, or volume. Enter three values and solve for the fourth — essential for lab preparation, titrations, and dilutions.
mol
L
g/mol
Enter your values above to see the results.
Tips & Notes
- ✓Volume must be entered in liters — convert mL by dividing by 1000. A 250 mL flask requires 0.250 L. Entering mL directly gives a molarity 1,000 times too high.
- ✓For common lab chemicals, memorize key molar masses: NaCl = 58.44 g/mol; NaOH = 40.00 g/mol; HCl = 36.46 g/mol; H₂SO₄ = 98.08 g/mol; glucose = 180.16 g/mol; KNO₃ = 101.10 g/mol.
- ✓When preparing a standard solution, add solvent to the solute gradually and mix thoroughly before making up to the final volume in a volumetric flask — do not fill to mark first and then add solute.
- ✓Serial dilution is more accurate than single dilution for very low concentrations. Prepare a 10 mM stock from 1 M, then dilute again to 10 µM, rather than attempting a 100,000-fold dilution in one step.
- ✓Always record molarity with units (M or mol/L) and specify temperature — molarity is temperature-dependent because liquid volume changes with temperature, which is why analytical labs standardize at 20°C or 25°C.
Common Mistakes
- ✗Entering volume in mL instead of liters — 500 mL entered as 500 gives molarity 1,000× too low. Always divide mL by 1,000 before entering: 500 mL = 0.500 L.
- ✗Confusing molarity (mol/L solution) with molality (mol/kg solvent) — these are different concentration units. Molality is used for colligative property calculations; molarity is used for stoichiometry and dilutions.
- ✗Using the volume of solvent instead of volume of solution — molarity uses the total final solution volume, not the volume of water added. Fill to the volumetric flask mark, not add 1 L of water to the solute.
- ✗Rounding molar mass prematurely — using 58 g/mol for NaCl instead of 58.44 g/mol introduces 0.75% error. For 1 M in 1 L, that is a 4.4 g discrepancy in the weighed mass.
- ✗Treating molar concentration as mass concentration — 1 M NaCl is not 1 g/L. It is 58.44 g/L. Confusing these units is one of the most common errors when converting between molarity and mass per volume.
Molarity Calculator Overview
Molarity is the universal language of solution chemistry — every titration, dilution, stoichiometric calculation, and spectrophotometric assay in the laboratory begins with knowing the molar concentration of the reagents involved. The ability to calculate molarity in all directions (solving for mass, volume, or moles from any combination of knowns) is the most fundamental quantitative skill in analytical and preparative chemistry.
Core molarity formula:
M (mol/L) = n (mol) / V (L) | n = M × V | V = n / M
EX: Prepare 250 mL of 0.5 M NaOH. n = M × V = 0.5 × 0.250 = 0.125 mol. Mass = n × MW = 0.125 × 40.00 = 5.00 g. Weigh 5.00 g NaOH, dissolve in ~200 mL water, transfer to 250 mL volumetric flask, fill to mark.Molarity from mass using molar mass:
M = m (g) / (MW (g/mol) × V (L)) | Mass needed = M × MW × V
EX: 10.0 g of NaCl (MW = 58.44 g/mol) dissolved to make 500 mL solution → M = 10.0 / (58.44 × 0.500) = 10.0 / 29.22 = 0.342 MCommon molar masses for laboratory reagents:
| Compound | Formula | MW (g/mol) | Common Use | Typical Stock Concentration |
|---|---|---|---|---|
| Sodium chloride | NaCl | 58.44 | Buffer, osmolarity standard | 1 M, 5 M |
| Sodium hydroxide | NaOH | 40.00 | Base, titrant | 1 M, 10 M |
| Hydrochloric acid | HCl | 36.46 | Acid, titrant | 1 M, conc. ~12 M |
| Sulfuric acid | H₂SO₄ | 98.08 | Acid, dehydrating agent | Conc. ~18.4 M |
| Acetic acid | CH₃COOH | 60.05 | Buffer component | Conc. ~17.4 M |
| Glucose | C₆H₁₂O₆ | 180.16 | Cell culture media | 1 M |
| Potassium nitrate | KNO₃ | 101.10 | Electrolyte, fertilizer | 0.1 M, 1 M |
| EDTA | C₁₀H₁₆N₂O₈ | 292.24 | Chelating agent, buffer | 0.5 M |
| Unit | Definition | Relationship to Molarity | Use Case |
|---|---|---|---|
| M (molar) | mol/L solution | Reference unit | General lab chemistry |
| mM (millimolar) | mmol/L = 10⁻³ M | 1 mM = 0.001 M | Biochemistry, enzymology |
| µM (micromolar) | µmol/L = 10⁻⁶ M | 1 µM = 10⁻⁶ M | Pharmacology, trace analysis |
| nM (nanomolar) | nmol/L = 10⁻⁹ M | 1 nM = 10⁻⁹ M | High-sensitivity assays |
| Normality (N) | equivalents/L | N = M × n-factor | Acid-base titrations |
| % w/v | g/100 mL | M = (%×density×10)/MW | Concentrated reagents |
Frequently Asked Questions
Molarity (M) is the number of moles of solute per liter of solution: M = n/V. It differs from molality (mol/kg of solvent), mass concentration (g/L), and percent concentration (g/100 mL). Molarity is temperature-dependent because liquid volume expands with heat. Molality is temperature-independent (mass does not change with temperature), which is why it is preferred for boiling point elevation and freezing point depression calculations. For most laboratory stoichiometry and dilution calculations, molarity is the standard unit.
Mass needed (g) = Molarity (M) × Volume (L) × Molar Mass (g/mol). Example: prepare 500 mL of 0.1 M NaCl. Mass = 0.1 mol/L × 0.500 L × 58.44 g/mol = 2.922 g. Weigh 2.922 g of NaCl, dissolve in about 400 mL of water in a 500 mL volumetric flask, then add water to the 500 mL graduation mark. Never add solute to a full flask — the volume would exceed the target. For hygroscopic compounds (NaOH, CaCl₂), weigh quickly and account for absorbed moisture.
The dilution equation is C₁V₁ = C₂V₂, where C is concentration and V is volume. Example: dilute a 12 M HCl stock solution to 0.5 M in 1 L total. V₁ = C₂V₂/C₁ = (0.5 M × 1 L) / 12 M = 0.0417 L = 41.7 mL. Add 41.7 mL of concentrated HCl to about 800 mL of water in a 1 L volumetric flask, then fill to the 1 L mark. Always add acid to water, never water to concentrated acid — the exothermic mixing can cause dangerous splattering if done in reverse.
Normality (N) = Molarity × n-factor, where n-factor is the number of equivalents per mole. For acids: n-factor = number of H⁺ ions donated. 1 M H₂SO₄ = 2 N (donates 2 H⁺). 1 M HCl = 1 N. For bases: n-factor = number of OH⁻ ions donated. 1 M NaOH = 1 N; 1 M Ca(OH)₂ = 2 N. For oxidation-reduction reactions, n-factor = electrons transferred. Normality simplifies equivalence point calculations in acid-base and redox titrations, but molarity is more commonly used in modern chemistry.
For concentrated reagents sold by mass percent: M = (% × density × 1000) / (MW × 100). Example: concentrated H₂SO₄ is 98% pure, density 1.84 g/mL, MW 98.08 g/mol. M = (98 × 1.84 × 1000) / (98.08 × 100) = 180,320 / 9,808 = 18.4 M. For 37% HCl (density 1.19 g/mL, MW 36.46): M = (37 × 1.19 × 1000) / (36.46 × 100) = 12.1 M. These concentrated acid molarities are essential starting points for preparing dilute working solutions in the lab.
Molar concentration (same as molarity) directly links to spectrophotometric measurements through Beer-Lambert Law: A = ε × l × c, where A is absorbance, ε is the molar absorptivity (L/mol·cm), l is path length (cm), and c is molar concentration (mol/L). To find concentration from absorbance: c = A / (ε × l). Example: a solution of p-nitrophenol with ε = 18,300 L/mol·cm at 405 nm, measured in a 1 cm cuvette with A = 0.732. c = 0.732 / (18,300 × 1) = 4.0 × 10⁻⁵ M = 40 µM. Beer-Lambert Law is the basis of most quantitative spectrophotometric assays in chemistry and biochemistry.